CH101 Exam #1 Review
CH101 Exam #1 Review CH 101
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This 10 page Study Guide was uploaded by Lauren Dutch on Thursday September 8, 2016. The Study Guide belongs to CH 101 at University of Alabama - Tuscaloosa taught by Dave Nikles in Fall 2016. Since its upload, it has received 191 views. For similar materials see General Chemistry in Chemistry at University of Alabama - Tuscaloosa.
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Date Created: 09/08/16
I. Modern Atomic Theory and Its Laws A. Law of Conservation of Mass- In a chemical reaction, matter is neither created nor destroyed. B. Law of Definite Proportions- All samples of a given compound, regardless of how they were made, have the same proportions of their constituent elements. C. Law of Multiple Proportions- When two elements A and B form two different compounds, the masses of element B that combine with 1g of element A can be expressed as a ratio of a small whole numbers. II. Discovery of Electron A. J.J. Thomson’s Cathode Rays 1. Cathode rays travel through a glass cathode ray tube away from the negative cathode towards the positive anode. 2. Discovery of a negatively charged, low mass particle present in all atoms called the electron B. Millikan’s Oil Drop Experiment 1. Calculated charge on oil droplets falling in and electric field 2. Discovered that the charge on an electron is -1.60x10^-19 C. Ernest Rutherford’s Gold Foil Experiment 1. Bombarded thin gold foil with positive particles a) Most passed through but some were deflected b) Denounced the plum pudding model that an atom is a positive sphere with negative electrons in it 2. Discovered the structure of the atom and explained it with the nuclear theory a) Most of the atom’s mass and all of its positive charge is contained in the nucleus b) Most of the volume is empty space, where the tiny negative electrons are found c) There are the same number of protons in nucleus and electrons outside the nucleus so the atom is neutrally charged III. Nuclear Model A. Almost all mass resides in the nucleus 1. Number of protons determines the atomic number 2. Protons and neutrons determine atomic mass (amu) 3. Isotopes have same atomic number (number of protons) but different atomic masses (different number of neutrons) a) Carbon-14 is an isotope of carbon. There are still 6 protons but there are two extra neutrons. 4. An element’s atomic mass is calculated by adding together all the different isotopes masses times their abundance a) Chlorine has two naturally occurring isotopes. Chlorine-35 has a mass of 34.97 amu and an abundance of 75.77%. Chlorine-37 has a mass of 36.97 amu and an abundance of 24.23%. (1) Atomic mass= 0.7577(34.97amu) + 0.2423(36.97amu) = 35.45amu B. Mass spectrometry calculates the mass of isotopes and their abundances by separating particles according to mass 1. Creates a mass spectrum a) X axis indicates mass b) Y axis indicates relative abundance c) Mass spectrum of chlorine CH101 Chapter 2 IV. Scientific Measurements A. Accuracy- how close you are to a true or accepted value B. Precision- reproducibility of a measurement C. Resolution- distinguishing parts of an object D. Significant figures 1. Multiplication and division- answer with the lowest number of sig figs from the problem a) Example: 2.5 * 3.42=8.6 2. Addition and subtraction- answer with lowest number of decimal places from the problem a) 34.56-5.4=29.2 E. Must have a value AND a unit 1. SI unit of length is the meter (m) 2. Prefixes a) Macroscale- human eye can see unaided b) Microscale- need an optical microscope to see c) Nanoscale- need an electron microscope to see V. Converting between moles and atoms A. Avogadro’s number: 1 mol=6.022x10^23 atoms B. 1 mole of an element equals the amu of that element in grams C. Example: Calcuate the number of carbon atoms in a 0.035 gram pencil lead. 0.035 g C (1 mol C / 12.01 g C) (6.022x10^23 atoms C / 1 mol C) = 1.75 atoms C Grams and moles cancel out, leaving atoms CH101 Chapter 3 The quantum-mechanical model of the atom explains the strange behavior electrons. I. Nature of Light A. Electromagnetic radiation 1. Photons travelling through space at the speed of light, c = 3x10^8 m/s 2. Oscillating electric and magnetic fields that resemble a sine curve a) Amplitude is the vertical height of a crest of depth of a trough from the x axis (1) Determines intensity/brightness- greater amplitude = greater intensity b) Wavelength is the distance between two crests (1) Symbol λ (2) Determines color (3) Measured in meters, micrometers, or nanometers c) Frequency is the number of cycles in a given period of time (1) Symbol v (2) Inversely related to wavelength (3) Measured in cycle/s or s^-1 or Hertz (Hz) d) Relationship between frequency and wavelength V = c / λ(c = speed of light) B. Electromagnetic Spectrum 1. Contains all the wavelengths of electromagnetic radiation 2. Lower frequency = longer wavelength = less energy Higher frequency = shorter wavelength = more energy C. Photoelectric Effect 1. Einstein discovered that light is quantized, meaning that light energy comes in packets, or photons. a) Amount of energy in packet: E = hv Where h is Plank’s constant (6.626x10^-34 J*s) and v is frequency b) Since v = c / λ the equation can also read E = hc / λ 2. When electromagnetic radiation of the proper wavelength strikes a material, photoelectrons are ejected from the sample. a) E is the minimum energy needed to generate photoelectrons (1) Must overcome binding energy, the energy with which the electron is bound to the metal (a) Low frequency, long wavelength light does not eject photoelectrons (b) High frequency, short wavelength light does eject photoelectrons (c) Threshold frequency condition: hv =Ф (Ф is binding energy and hv is energy of photon) (d) Kinetic energy is the difference between the energy of the photon and the binding energy KE = hv – Ф II. Atomic Spectroscopy A. Emission spectra is a tool to identify the type and amount of elements present 1. A hydrogen lamp shines through a prism to create a line spectrum 2. Each element has a unique emission spectrum B. An absorption spectrum can also be used to identify elements but instead of having bright lines on a dark background, it has dark lines on a bright background C. Bohr Model 1. When an electron transitions from one stationary state to another, it emits a photon with energy equal to the energy difference of the two states a) States that are close together produce lower energy, longer wavelength light while states that are further apart produce higher energy, shorter wavelength light. 2. Quantum Leaps a) Electronic excitation caused by visible or UV light b) Radiation is when energy is given off as an electron falls back to a lower energy level (1) If an electron moves up, it will eventually fall back down to produce light c) Energy from excitation or radiation can be calculated E = Efinal – Einitial En= -2.179x10^-18 / n^2 J d) Rydberg equation is used to find the wavelength produced when an electron moves states 1 / λ = R (1 / n^2 – 1 / m^2) R = Rydberg constant = 1.097 x 10^7 e) Flame tests can also identify elements when a sample is placed in a Bunsen burner flame and the resulting color is observed III. Wave Nature of Matter A. De Broglie Wavelength 1. Calculate wavelength of a moving electron Λ = h / mv M is mass, v is velocity B. Heisenberg’s Uncertainty Principle 1. We cannot observe the wave nature and the particle nature of electrons at the same time IV. Quantum Mechanics A. All matter consists of atoms in constant motion B. Wave functions are electron orbitals C. Schrodinger wave equation D. Quantum numbers 1. Variables in wave functions of Schrodinger wave equation 2. Each orbital can only hold 2 electrons 3. Principal = n a) All positive integers besides 0 b) By row on periodic table 4. Azimuthal = L a) Determines shape of orbital b) Up to n – 1 c) Example: if n = 3, L can be 0, 1, or 2 5. Magnetic = m L a) Determines direction of orbital b) Can be any integer between –L and +L 6. Spin = m s a) Determines direction of electrons with respect to external magnetic field b) -1/2 or +1/2 c) Spins cancel out if two electrons are in orbital V. Shapes of orbitals A. S orbital 1. L = 0, m L 0 2. Spherical B. P orbital 1. L = 1, m L -1, 0, 1 a) Because m hLs 3 options, there are 3 degenerate orbitals in this level, meaning the orbitals have the same energy 2. Dumbbells 3. Node at origin of dumbbell shape; no probability of finding an electron at a node C. D orbital 1. L = 2, m L -2, -1, 0, 1, 2 a) Five degenerate orbitals 2. Node at origin 3. Four leaf clover D. When orbitals interact, they can either be in phase (same sign on wave function) or out of phase (opposite signs on wave function)
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